Class 12 Chemistry · NCERT chapter 3 · 8 topics · free to read
Electrochemical Cells
Electrochemical cells convert chemical energy to electrical energy (galvanic/voltaic) or electrical energy to chemical energy (electrolytic). A galvanic cell has two half-cells connected by a salt bridge. Cell notation follows the convention: anode (oxidation) on the left, cathode (reduction) on the right. The Daniel cell is the classic example: Zn|Zn²⁺||Cu²⁺|Cu.
A galvanic (voltaic) cell converts the energy of a spontaneous redox reaction into electrical energy. It consists of two half-cells: one where oxidation occurs (anode) and one where reduction occurs (cathode), connected externally by a wire and internally by a salt bridge.
A galvanic cell (also called voltaic cell) is a device that converts chemical energy into electrical energy using a spontaneous redox reaction. The key idea: separate the oxidation and reduction half-reactions into two different containers (half-cells), and force the electrons to travel through an external wire — that electron flow is your electric current.
Anode = electrode where oxidation occurs (electrons are released). Anode is negative in a galvanic cell
Cathode = electrode where reduction occurs (electrons are consumed). Cathode is positive in a galvanic cell
Electrons flow from anode → cathode through the external circuit
Conventional current flows from cathode → anode (opposite to electron flow)
The cell produces a positive EMF ( E_ cell > 0 ) because the reaction is spontaneous ( G < 0 )
The Daniel Cell
The Daniel cell uses a zinc anode dipping in ZnSO₄ solution and a copper cathode dipping in CuSO₄ solution. Zinc gets oxidized (dissolves), copper ions get reduced (deposit). The standard EMF is 1.10 V.
The Daniel cell is the most important galvanic cell for NEET. It consists of: - Anode half-cell: Zn rod in ZnSO_4 solution (1 M) - Cathode half-cell: Cu rod in CuSO_4 solution (1 M) - Salt bridge: connects the two solutions (usually KCl or KNO₃ in agar)
A salt bridge is a U-tube containing an inert electrolyte (KCl, KNO₃) in agar-agar. It maintains electrical neutrality in both half-cells by allowing ion migration, and completes the internal circuit.
Without a salt bridge, a galvanic cell stops working within seconds. Why? As Zn dissolves, the anode solution accumulates excess Zn^ 2+ (positive charge builds up). As Cu^ 2+ deposits, the cathode solution loses positive ions (negative charge builds up). This charge imbalance quickly stops the reaction.
Made of a U-tube with an inert electrolyte (KCl, KNO₃, NH_4NO_3 ) in agar-agar gel
Functions: (1) Maintains electrical neutrality in both half-cells, (2) Completes the internal circuit, (3) Minimizes liquid junction potential
The electrolyte must be inert — it should NOT react with either solution
KCl is most commonly used because K^+ and Cl^- have nearly equal transport numbers (similar migration speeds)
If the salt bridge is removed, the cell voltage drops to zero almost immediately
Quick recall
Why can't you use KCl as a salt bridge for a cell containing Ag^+ ions?
Because Cl^- from KCl would react with Ag^+ to form insoluble AgCl precipitate. Use KNO_3 instead — NO_3^- doesn't precipitate with most metal ions.
Cell Notation
Cell notation (cell representation) is a compact way to describe a galvanic cell. The convention: anode on left, cathode on right, single vertical line for phase boundary, double vertical line for salt bridge.
Instead of drawing the full cell diagram, chemists use cell notation (also called cell representation). It follows a strict convention that NEET loves to test.
Single vertical line | = phase boundary (solid|solution or solution|gas, etc.)
Double vertical line || = salt bridge
Concentrations are written in parentheses after the species
For a gas electrode: Pt(s) | H_2(g, 1 atm) | H^+(aq, 1M) — Pt acts as the inert electrode
Galvanic vs Electrolytic
Galvanic cells use spontaneous reactions to generate electricity; electrolytic cells use external electricity to drive non-spontaneous reactions. Electrode polarity is reversed between the two types.
These are the two fundamental types of electrochemical cells. A galvanic cell is a battery — it generates electricity from a spontaneous reaction. An electrolytic cell is the opposite — it uses external electricity to force a non-spontaneous reaction.
Chemical energy → Electrical energy
Electrical energy → Chemical energy
Spontaneous reaction ( G < 0 )
Non-spontaneous reaction ( G > 0 )
E_ cell > 0 (positive EMF)
E_ cell < 0 (needs external EMF)
Two separate containers + salt bridge
Both electrodes in the SAME container
Anode is negative (−)
Anode is positive (+)
Cathode is positive (+)
Cathode is negative (−)
Example: Daniel cell, dry cell
Example: electrolysis of water, electroplating
Quick recall
A rechargeable battery is being charged. Is it acting as a galvanic or electrolytic cell?
Electrolytic cell. During charging, external electricity forces the reverse (non-spontaneous) reaction. During discharging (use), it acts as a galvanic cell.
EMF & Standard Electrode Potential
The standard electrode potential (E°) measures the tendency of an electrode to gain electrons. E°cell = E°cathode − E°anode. The electrochemical series ranks elements by E° — metals with more negative E° are stronger reducing agents. A positive E°cell means the reaction is spontaneous.
E^ _ cell = E^ _ cathode - E^ _ anode
Standard Electrode Potential
The standard electrode potential (E°) is the potential difference developed between an electrode and its solution when all species are at standard conditions (1 M concentration, 1 atm pressure, 25°C). It is measured relative to the Standard Hydrogen Electrode (SHE).
We can't measure the absolute potential of a single electrode — we can only measure the difference between two electrodes. So we need a reference. That reference is the Standard Hydrogen Electrode (SHE), which is arbitrarily assigned a potential of 0.00 V.
Standard conditions: all ion concentrations = 1 M, gas pressures = 1 atm, temperature = 298 K (25°C)
SHE consists of a Pt electrode dipped in 1 M HCl with H_2 gas bubbled at 1 atm
Pt is used because it is inert (doesn't react) and catalytic (facilitates the half-reaction)
E^ > 0 means the electrode has a greater tendency to get reduced than SHE
E^ < 0 means the electrode has a lesser tendency to get reduced than SHE (it prefers to get oxidized)
Cell EMF Calculation
The EMF (electromotive force) of a cell equals the standard reduction potential of the cathode minus that of the anode. A positive EMF means the cell reaction is spontaneous.
The EMF (electromotive force) of a galvanic cell is the potential difference between the two electrodes when no current is drawn. It equals the maximum work the cell can do per coulomb of charge.
E^ _ cell = E^ _ cathode - E^ _ anode
If E^ _ cell > 0 → reaction is spontaneous ( G^ < 0 )
If E^ _ cell < 0 → reaction is non-spontaneous ( G^ > 0 )
If E^ _ cell = 0 → system is at equilibrium
EMF is measured using a potentiometer (not a voltmeter) because a voltmeter draws current and gives a lower reading
EMF is an intensive property — independent of the size of the cell or the amount of reactants
Quick recall
Given E^ (Ag^+/Ag) = +0.80 V and E^ (Mg^ 2+ /Mg) = -2.37 V, find E^ _ cell for Mg-Ag cell.
Mg is more active → anode. Ag is nobler → cathode. E^ _ cell = +0.80 - (-2.37) = +3.17 V. The large positive value means a highly spontaneous reaction.
Electrochemical Series
The electrochemical series lists standard reduction potentials in increasing order. Elements at the top (most negative E°) are strongest reducing agents; elements at the bottom (most positive E°) are strongest oxidizing agents. Used to predict spontaneity, reactivity, and displacement reactions.
The electrochemical series (also called the activity series or EMF series) arranges half-reactions in order of their standard reduction potentials. It's the master reference table for predicting which reactions will occur spontaneously.
Top of series (most negative E^ ): Li, K, Ca, Na, Mg, Al, Zn, Fe — strongest reducing agents (easily oxidized)
Bottom of series (most positive E^ ): Au , Pt , F_2 — strongest oxidizing agents (easily reduced)
A metal higher in the series can displace a metal lower in the series from its salt solution
Any metal above hydrogen can displace H_2 from dilute acids
SHE ( E^ = 0 ) is the dividing line — metals above it are reactive; metals below it are noble
Quick recall
Will iron displace copper from CuSO_4 solution?
Yes! Fe ( E^ = -0.44 V) is above Cu ( E^ = +0.34 V) in the series. Fe(s) + CuSO_4(aq) arrow FeSO_4(aq) + Cu(s) . That's why dipping an iron nail in blue CuSO_4 turns it brownish (Cu coating).
Nernst Equation
The Nernst equation gives the cell potential under non-standard conditions by accounting for ion concentrations. At 25°C: E = E° − (0.0591/n) log Q. When Q = K (equilibrium), E = 0. Concentration cells have E° = 0 but E ≠ 0 due to concentration difference.
E_ cell = E^ _ cell - 0.0591 n Q (at 298 K)
Nernst Equation
The Nernst equation relates cell EMF to standard EMF and the reaction quotient Q. It shows how changing concentrations or pressures shifts the cell potential from its standard value.
The standard EMF ( E^ ) assumes all concentrations are 1 M and all pressures are 1 atm. In reality, concentrations vary. The Nernst equation tells you the actual EMF under any conditions.
R = 8.314 J mol ^ -1 K ^ -1 , T = 298 K, F = 96485 C mol ^ -1
n = number of moles of electrons transferred in the balanced cell reaction
Q = reaction quotient = [products] [reactants] (same expression as equilibrium constant, but at current concentrations)
The factor 2.303RT F = 2.303 8.314 298 96485 = 0.0591 V at 298 K
Pure solids and pure liquids do NOT appear in the Q expression
Nernst Equation for a Half-Cell
The Nernst equation applies to individual half-cells too. For a metal/metal-ion electrode: E = E° − (0.0591/n) log(1/[Mⁿ⁺]).
The Nernst equation works for individual half-cells (single electrodes), not just full cells. For a reduction half-reaction M^ n+ (aq) + ne^- arrow M(s) :
E_ M^ n+ /M = E^ _ M^ n+ /M - 0.0591 n 1 [M^ n+ ]
E_ M^ n+ /M = E^ _ M^ n+ /M + 0.0591 n [M^ n+ ]
As [M^ n+ ] increases → electrode potential increases (more positive) → stronger tendency for reduction
As [M^ n+ ] decreases → electrode potential decreases (more negative) → stronger tendency for oxidation
At [M^ n+ ] = 1 M → 1 = 0 → E = E^ (standard conditions recovered)
For Cu^ 2+ /Cu electrode ( E^ = +0.34 V), what happens to E if [Cu^ 2+ ] is reduced from 1 M to 0.01 M?
E = 0.34 + 0.0591 2 (0.01) = 0.34 + 0.02955 (-2) = 0.34 - 0.0591 = 0.281 V. Diluting the Cu^ 2+ solution decreases the reduction potential.
Concentration Cells
A concentration cell has identical electrodes in solutions of different concentrations. E° = 0 (same electrode material), but E ≠ 0 because of the concentration difference. EMF drives ions from high concentration to low concentration.
A concentration cell uses the same metal and the same ions in both half-cells — the ONLY difference is the concentration of the electrolyte. Since both electrodes are identical, E^ _ cell = 0 . But the Nernst equation still gives a non-zero EMF!
E_ cell = 0 - 0.0591 n C_1 C_2 = 0.0591 n C_2 C_1
where C_2 > C_1 (higher concentration at cathode). The cell works to equalize concentrations — dilute side gets oxidized (anode), concentrated side gets reduced (cathode).
The electrode in the dilute solution is the anode (oxidation: metal dissolves, adding ions)
The electrode in the concentrated solution is the cathode (reduction: ions deposit as metal)
The cell operates until both concentrations become equal → then E = 0 and cell stops
EMF depends on the ratio C_2/C_1 , not the absolute values
A 10-fold concentration ratio gives E = 0.0591 n 1 = 0.0591 n V
Quick recall
A Cu concentration cell has [Cu^ 2+ ] = 0.001 M and 0.1 M. Find the EMF.
E = 0.0591 2 0.1 0.001 = 0.02955 (100) = 0.02955 2 = 0.0591 V.
Nernst Equation at Equilibrium
At equilibrium, the cell potential drops to zero (E = 0), and Q becomes equal to the equilibrium constant K. This links electrochemistry to thermodynamics: log K = nE°/0.0591.
A galvanic cell produces current until the reaction reaches equilibrium. At equilibrium, the driving force for the reaction is zero — so the cell voltage drops to zero: E_ cell = 0 and Q = K .
0 = E^ _ cell - 0.0591 n K
K = n E^ _ cell 0.0591
At equilibrium: E = 0 , Q = K , G = 0
If E^ > 0 → K > 0 → K > 1 → products are favored
If E^ < 0 → K < 0 → K < 1 → reactants are favored
A large positive E^ means an astronomically large K — the reaction goes essentially to completion
For the Daniel cell: K = 2 1.10 0.0591 = 37.2 → K = 10^ 37.2 — practically irreversible!
Gibbs Energy & Cell Potential
The relationship ΔG° = −nFE° connects thermodynamics to electrochemistry. A positive E°cell means negative ΔG° (spontaneous). The equilibrium constant K can be found from E° via log K = nE°/0.0591 at 25°C. Maximum useful work from a cell equals −ΔG = nFE.
G^ = -nFE^ _ cell
ΔG° = −nFE°
Gibbs free energy change and cell EMF are directly related. A spontaneous cell reaction (positive E°) has a negative ΔG°. The maximum electrical work a cell can do equals the decrease in Gibbs energy.
The work done by a galvanic cell is electrical work: moving n moles of electrons through a potential difference E . Each mole of electrons carries a charge of F (Faraday constant = 96485 C/mol). So total charge = nF , and electrical work = charge potential = nFE .
G = -nFE_ cell and G^ = -nFE^ _ cell
n = number of moles of electrons transferred in the balanced equation
F = Faraday constant = 96485 C mol ^ -1 96500 C mol ^ -1
The negative sign makes the equation consistent: spontaneous reaction → E^ > 0 → G^ < 0
| G^ | = maximum useful work obtainable from the cell (at reversible conditions)
If E^ _ cell > 0 → G^ < 0 → spontaneous
If E^ _ cell < 0 → G^ > 0 → non-spontaneous
Quick recall
Calculate G^ for the Daniel cell ( E^ = 1.10 V, n = 2 ).
G^ = -nFE^ = -2 96485 1.10 = -212267 J = -212.3 kJ mol ^ -1 . The large negative value confirms the reaction is highly spontaneous.
Finding K from E°
The standard EMF, Gibbs energy, and equilibrium constant are all interconnected. From E°cell, you can calculate both ΔG° and K. The key relation: log K = nE°cell/0.0591 at 25°C.
Three fundamental thermodynamic/electrochemical quantities are linked in a triangle. Knowing any one gives you the other two.
G^ = -nFE^ = -RT K
K = nE^ 0.0591 (at 298 K)
From E^ to G^ : G^ = -nFE^
From E^ to K : K = nE^ /0.0591
From G^ to K : G^ = -2.303 RT K
All three quantities tell the SAME story about spontaneity — just in different languages
K = 2 0.295 0.0591 = 0.59 0.0591 = 9.98 10 . So K = 10^ 10 . This means products are overwhelmingly favored.
Conductance & Conductivity
Conductance (G) is the reciprocal of resistance. Specific conductance (κ) measures the conducting ability of a solution per unit volume. Molar conductivity (Λm) = κ/c measures conductance per mole of electrolyte. κ decreases with dilution but Λm increases with dilution for both strong and weak electrolytes.
_m = 1000 c (c in mol/L, in S cm^ -1 )
Conductance & Resistance
Resistance (R) opposes current flow; conductance (G = 1/R) measures how easily current flows. For electrolytic solutions, R depends on concentration, temperature, and the nature of the electrolyte.
An electrolytic solution conducts electricity through the movement of ions. How well it conducts depends on the number, charge, and mobility of the ions. We measure this as conductance — the reciprocal of resistance.
G = 1 R Unit: siemens (S) or ^ -1 or mho
R = l A G = A l
R = resistance ( ), = resistivity ( cm), l = distance between electrodes (cm), A = area of electrodes (cm ^2 )
G = conductance (S), = conductivity or specific conductance (S cm ^ -1 )
= 1 = conductivity = conductance of a solution confined between two electrodes 1 cm apart with 1 cm ^2 area
Conductance is measured using a Wheatstone bridge with an AC source (to prevent electrolysis) — called a conductivity bridge
l A is called the cell constant ( G^* ), unit: cm ^ -1 . So = G G^* = G l A
Specific Conductance (κ)
Specific conductance (κ, kappa) is the conductance of a unit cube of solution. It depends on concentration: for electrolytes, κ increases with concentration because more ions are available per unit volume.
Specific conductance ( ) is the conductance of a solution between two parallel electrodes of 1 cm ^2 area placed 1 cm apart. It tells you the intrinsic conducting ability of the solution.
= G^* R = l A R Unit: S cm^ -1 or S m^ -1
increases with increasing concentration (more ions per cm ^3 ) — but only up to a point
At very high concentrations, interionic attractions reduce ion mobility → may decrease
increases with increasing temperature (ions move faster, viscosity decreases)
Strong electrolytes have higher than weak electrolytes at the same concentration
SI unit: S m ^ -1 . CGS unit: S cm ^ -1 . Conversion: 1 S cm ^ -1 = 100 S m ^ -1
Molar Conductivity (Λm)
Molar conductivity Λm = κ/c is the conductance of all ions produced by one mole of electrolyte dissolved in a given volume. It increases with dilution and approaches a limiting value Λ°m at infinite dilution.
Molar conductivity ( _m ) is the conducting power of all the ions produced by dissolving one mole of electrolyte in a given volume of solution. It answers: 'If I take one mole of electrolyte, how well does it conduct?'
_m = c where c = molar concentration
If is in S cm ^ -1 and c is in mol L ^ -1 (molarity):
_m = 1000 c Unit: S cm^2 mol^ -1
_m increases with dilution for ALL electrolytes — because interionic interactions decrease
For strong electrolytes: _m increases slowly and reaches ^ _m (limiting molar conductivity) at moderate dilution
For weak electrolytes: _m increases steeply near infinite dilution because degree of dissociation ( ) increases dramatically
^ _m = molar conductivity at infinite dilution (when the solution is so dilute that ions behave independently)
For weak electrolytes: = _m ^ _m (degree of dissociation from conductivity data)
Kohlrausch's Law
Kohlrausch's law of independent migration of ions states that the limiting molar conductivity of an electrolyte equals the sum of individual ion conductivities: Λ°m = ν+λ°+ + ν−λ°−. This is critical for finding Λ°m of weak electrolytes (like CH₃COOH) that cannot be measured directly, and for calculating solubility of sparingly soluble salts.
^ _m = _+ ^ _+ + _- ^ _-
Kohlrausch's Law
At infinite dilution, each ion migrates independently of its co-ion. The limiting molar conductivity of an electrolyte is the sum of the limiting ionic conductivities of its cation and anion, each multiplied by the number of ions produced.
At infinite dilution, ions are so far apart that they don't interact with each other at all. Each ion moves independently, carrying its own share of the current. This is Kohlrausch's law of independent migration of ions.
^ _m = _+ ^ _+ + _- ^ _-
^ _m = limiting molar conductivity of the electrolyte (at infinite dilution)
^ _+ and ^ _- = limiting ionic conductivities of the cation and anion
_+ and _- = number of cations and anions per formula unit
Weak electrolytes don't fully dissociate, so Λ°m can't be found by extrapolation. Kohlrausch's law lets us calculate it from the Λ°m values of strong electrolytes through algebraic addition and subtraction.
For strong electrolytes, you can find ^ _m by plotting _m vs c and extrapolating to c = 0 . But for weak electrolytes (like CH_3COOH ), _m shoots up steeply near infinite dilution — extrapolation is impossible. Kohlrausch's law provides the workaround.
Strategy: Express ^ _m of the weak electrolyte as a combination of strong electrolytes whose ^ _m values are known.
Choose strong electrolytes that together contain ALL the ions of the weak electrolyte
Add/subtract to cancel out unwanted ions
The final expression should give: ^ (cation of weak electrolyte) + ^ (anion of weak electrolyte)
This is the most frequently tested application of Kohlrausch's law in NEET
For sparingly soluble salts, the saturated solution is extremely dilute, so Λm ≈ Λ°m. Using κ of the saturated solution and Λ°m (from Kohlrausch's law), you can find the molar solubility.
For a sparingly soluble salt (like AgCl, BaSO_4 ), the saturated solution is so dilute that _m ^ _m . We can exploit this to find solubility from conductivity measurements.
Steps: 1. Measure of the saturated solution 2. Subtract of pure water: _ salt = _ solution - _ water 3. Calculate ^ _m using Kohlrausch's law 4. Find molar solubility: c = _ salt 1000 ^ _m mol/L
c = _ salt 1000 ^ _m mol L^ -1
This works only for sparingly soluble salts where the solution is effectively at infinite dilution
You MUST subtract the conductivity of pure water — even 'pure' water has 5 10^ -5 S cm ^ -1
Once you have molar solubility c , you can find K_ sp = c^2 (for 1:1 salt like AgCl)
For K_ sp of AB_2 type (like PbCl_2 ): K_ sp = 4c^3
Quick recall
The conductivity of a saturated AgCl solution (after subtracting water) is 1.86 10^ -6 S cm ^ -1 . If ^ _m(AgCl) = 138.3 S cm ^2 mol ^ -1 , find the molar solubility.
c = 1.86 10^ -6 1000 138.3 = 1.345 10^ -5 mol L ^ -1 .
Degree of Dissociation from Conductivity
For weak electrolytes, the ratio of molar conductivity at a given concentration to the limiting molar conductivity gives the degree of dissociation. Combined with Ostwald's dilution law, it also gives Ka.
For a weak electrolyte like acetic acid, only a fraction of the molecules dissociate into ions. This fraction is the degree of dissociation ( ), and it can be determined from conductivity data.
= _m ^ _m
Combining with Ostwald's dilution law for a weak acid HA:
K_a = c ^2 1 - = c _m^2 ^ _m( ^ _m - _m)
ranges from 0 (no dissociation) to 1 (complete dissociation)
At infinite dilution, = 1 for ALL electrolytes (even weak ones) → _m = ^ _m
For strong electrolytes, 1 at all concentrations (so this formula is mainly useful for weak electrolytes)
Electrolysis uses electricity to drive non-spontaneous reactions. Faraday's first law: m = ZIt (mass deposited ∝ charge). Faraday's second law: masses of different substances deposited by the same charge are proportional to their equivalent weights. One faraday (96485 C) deposits one gram-equivalent of any substance.
m = M I t n F = Z I t
Electrolysis
Electrolysis is the process of using electrical energy to decompose or drive a non-spontaneous chemical reaction. The electrolyte (molten or aqueous) provides ions that migrate to electrodes under the applied voltage.
In electrolysis, an external battery forces a non-spontaneous reaction to occur. The setup: two electrodes dipped in an electrolyte (molten salt or aqueous solution), connected to a DC source. Cations migrate to the cathode (reduction), anions migrate to the anode (oxidation).
Cations (positive ions) move toward the cathode (negative electrode) → get reduced
Anions (negative ions) move toward the anode (positive electrode) → get oxidized
In an electrolytic cell: anode is positive (+), cathode is negative (−) — opposite to galvanic cell polarity
The minimum voltage needed to start electrolysis is called the decomposition potential
If multiple ions are present, the ion that gets reduced/oxidized preferentially depends on its E^ value
Products of Electrolysis
The products of electrolysis depend on the electrolyte (molten vs aqueous), the electrode material (inert vs active), and the standard electrode potentials of competing reactions.
Predicting electrolysis products is a favorite NEET question. The key: when multiple species can be oxidized/reduced, the one requiring the least energy wins.
Aqueous NaCl (concentrated): Cathode: H_2 gas (water reduced); Anode: Cl_2 gas (Cl ^- oxidized due to overpotential effects)
Dilute H_2SO_4 : Cathode: H_2 gas; Anode: O_2 gas (water oxidized at both electrodes — it's water electrolysis!)
Aqueous CuSO_4 with Pt electrodes: Cathode: Cu metal deposits; Anode: O_2 gas
Aqueous CuSO_4 with Cu electrodes: Cathode: Cu deposits; Anode: Cu dissolves — copper is transferred from anode to cathode (copper refining!)
Faraday's First Law
The mass of substance deposited or dissolved at an electrode during electrolysis is directly proportional to the quantity of electricity (charge) passed through the solution. m = ZIt, where Z is the electrochemical equivalent.
Michael Faraday discovered the quantitative laws of electrolysis in 1834. His first law says: the amount of chemical change (mass deposited/dissolved) is directly proportional to the quantity of electricity passed.
m = Z I t = Z Q
Z = M nF (electrochemical equivalent)
m = M I t n F = M Q n F
m = mass deposited (g), Q = charge passed = I t (coulombs)
I = current (amperes), t = time (seconds)
Z = electrochemical equivalent = mass deposited per coulomb (g/C)
M = molar mass (g/mol), n = number of electrons in the half-reaction, F = 96485 C/mol
1 Faraday = 96485 C = charge carried by 1 mole of electrons = deposits 1 gram-equivalent of any substance
Quick recall
How much copper deposits when 2 A current is passed through CuSO_4 for 1 hour? ( M_ Cu = 63.5 , n = 2 )
Q = I t = 2 3600 = 7200 C. m = 63.5 7200 2 96485 = 457200 192970 = 2.37 g.
Faraday's Second Law
When the same quantity of electricity is passed through different electrolytes connected in series, the masses of substances deposited are proportional to their equivalent weights (M/n).
Faraday's second law compares what happens in different electrolytic cells when the SAME charge flows through all of them (connected in series).
m_1 m_2 = E_1 E_2 = M_1/n_1 M_2/n_2
E = M/n = equivalent weight = molar mass divided by the number of electrons transferred
Same charge through series cells deposits same number of equivalents of each substance
1 Faraday deposits exactly 1 gram-equivalent of any substance: 107.87/1 = 107.87 g Ag, 63.5/2 = 31.75 g Cu, 27/3 = 9 g Al
This law is the basis for the copper voltameter (used to measure charge)
Quick recall
Two cells (AgNO₃ and CuSO₄) are connected in series. If 1.08 g Ag is deposited, how much Cu is deposited?
m_ Cu m_ Ag = E_ Cu E_ Ag = 31.75 107.87 . So m_ Cu = 1.08 31.75 107.87 = 0.318 g.
Corrosion & Batteries
Corrosion (rusting) is an electrochemical process where metal is oxidized by environmental agents. Rust forms through microscopic galvanic cells on the iron surface. Prevention: galvanization, cathodic protection, coating. Batteries are practical galvanic cells — primary (non-rechargeable: dry cell, mercury), secondary (rechargeable: lead-acid, Li-ion), and fuel cells (H₂-O₂).
Corrosion is the gradual destruction of metals by electrochemical reactions with the environment. Iron rusting is the most common example — it involves formation of tiny galvanic cells on the metal surface where iron is the anode (oxidized) and impurities or surface variations act as cathodes.
Rusting is NOT simply 'iron reacting with oxygen and water'. It's an electrochemical process. Different regions on the iron surface act as tiny anodes and cathodes, forming microscopic galvanic cells. Water droplets containing dissolved O_2 and CO_2 (which provide H^+ ) act as the electrolyte.
At the anode (iron surface — often at scratches, grain boundaries, or stressed areas):
Fe(s) arrow Fe^ 2+ (aq) + 2e^- (oxidation)
At the cathode (nearby region with access to O_2 ):
Fe^ 2+ is further oxidized to Fe^ 3+ , which forms hydrated iron(III) oxide — the familiar reddish-brown rust: Fe_2O_3 xH_2O .
Rusting requires BOTH water (moisture) and oxygen — neither alone causes rust
Dissolved electrolytes (like salt) ACCELERATE rusting — that's why coastal areas have more corrosion
Rust is not protective — it flakes off, exposing fresh iron (unlike Al_2O_3 on aluminum, which protects)
Acidic conditions speed up rusting ( H^+ participates in the cathode reaction)
Prevention of Corrosion
Corrosion can be prevented by barrier methods (painting, oiling, electroplating), sacrificial anode/cathodic protection (attaching a more reactive metal), galvanization (zinc coating), and using anti-rust solutions.
Since corrosion requires moisture, oxygen, and an electrolyte, prevention strategies block one or more of these factors — or use electrochemistry itself as a defense.
Electroplating: Coating with another metal (Cr, Ni, Sn) — protects as long as the coating is intact
Galvanization: Coating iron with zinc. Even if the zinc coating is scratched, zinc (more reactive) acts as the sacrificial anode and iron is protected
Cathodic protection (sacrificial anode): Attach a block of more reactive metal (Zn, Mg) to the iron structure. The reactive metal corrodes instead of iron. Used for underground pipelines, ship hulls
Alloying: Stainless steel (Fe + Cr + Ni) — chromium forms a protective Cr_2O_3 layer
Anti-rust solutions: Alkaline phosphate or chromate solutions form a protective film on iron
Primary Batteries
Primary batteries are non-rechargeable galvanic cells. The dry cell (Leclanché cell) uses Zn/MnO₂ and produces 1.5 V. The mercury cell uses Zn/HgO and produces a very stable 1.35 V. Once the reactants are exhausted, the battery is dead.
Primary batteries cannot be recharged — the cell reaction is irreversible (or practically so). Once the reactants are consumed, you throw the battery away.
Dry cell (Leclanché cell): Anode: Zn casing. Cathode: carbon rod surrounded by MnO_2 + carbon black. Electrolyte: NH_4Cl + ZnCl_2 paste. EMF ≈ 1.5 V. Used in torches, remotes, clocks
Secondary batteries are rechargeable — the cell reaction can be reversed by applying external voltage. The lead-acid battery (car battery) and nickel-cadmium battery are classic examples. The lead-acid battery provides 2 V per cell and is used in automobiles.
A secondary battery (rechargeable/storage cell) can be recharged by passing current in the reverse direction, converting electrical energy back into chemical energy. The reaction reverses, regenerating the original reactants.
Each cell gives 2 V; a car battery has 6 cells in series → 12 V
Electrolyte: dilute H_2SO_4 (specific gravity ≈ 1.28 when fully charged)
During discharge: H_2SO_4 is consumed → specific gravity decreases. Mechanic checks this with a hydrometer!
During charging: the reaction reverses — PbSO_4 converts back to Pb and PbO_2 , and H_2SO_4 is regenerated
Nickel-cadmium (Ni-Cd) battery: Longer life than lead-acid. Cd anode, NiO(OH) cathode, KOH electrolyte. EMF ≈ 1.2 V per cell
Fuel Cells
A fuel cell converts the chemical energy of a fuel (like H₂) directly into electrical energy without combustion. The H₂-O₂ fuel cell produces water as the only byproduct. Efficiency is much higher than combustion engines.
A fuel cell is a galvanic cell where the reactants are continuously supplied from outside. Unlike a regular battery, a fuel cell doesn't store energy — it converts fuel to electricity on the fly, as long as fuel keeps flowing.
Hydrogen-oxygen fuel cell (the most important for NEET):